Environ. Sci. Technol. 1993, 27, 304-310
Photochemical Reactions Involved in the Total Mineralization of 2,4-D by Fe3+/H,0,/UV Yunfu Sun and Joseph J. Plgnatello'
Department of Soil and Water, The Connecticut Agricultural Experiment Station, P.O. Box 1106, New Haven, Connecticut 06504
rn The photoenhancement of Fenton-type (Fe3++ HP02) oxidation of 2,4-dichlorophenoxyaceticacid in the near-UV is due to several Fe(II1)-sensitizedphotochemical reactions. Fe3+/HzOz/UVsystems were studied for potential use in waste treatment. Dechlorination and conversion of the first -40% of ring and carboxy carbon of 2,4-D to C02 are due mostly to hydroxyl radical (OH')reactions. Dark Fenton-generated OH' is supplemented by the following photoreactions that give OH' directly and/or provide in situ Fez+for the Fenton reaction: (i) photolysis of ferric ion, FeOH2+ Fez+ OH*;(ii) photodecarboxylation of a ferric-2,4-D complex, [RCO2-FeI2+ Fez++ COP+ W; and (iii) additional photochemical reaction(s) that may involve photolysis of a ferric-peroxy complex, Fe-OZH2+. The relative importance of each depends on [H202],except for photodecarboxylation, which is always minor. The remaining -60% of carbon mineralization occurs almost exclusively by photolysis/decarboxylation of Fe(II1) complexes of degradation intermediates, probably ring-opened products coordinated via carboxyl and hydroxyl groups; OH' plays no significant role during this stage. Oxalic acid was identified as one intermediate. Dioxygen is consumed in the overall reaction and accelerates carbon mineralization.
-
-
+
Introduction Reactions that generate hydroxyl radical (OH')in solution at low temperature have attracted interest for destruction of toxic organic compounds in wastewaters (I, 2). One way to generate OH' is by the well-known reactions of hydrogen peroxide with Fez+and Fe3+salts. The basic chemistry, as well as applications of Fe2+/HzO2and Fe3+/HzOZsystems for hazardous waste treatment, has been reviewed elsewhere (3). Briefly, ferrous ion combined with hydrogen peroxide ("Fenton's reagent") reacts stoichiometrically to give OH' Fez+-t HzOz Fe3+ OH- + OH' (1)
-
+
In the closely related ferric system, Fe3+acts as a catalyst for decompositionof peroxide to O2and HzO,during which "steady-state" concentrations of Fe2+(as a source of OH' via eq 1) are generated, as in the following steps: Fe3++ H202
-H+
Fe-O2H2+'
Fe3+ + HOP'
-
Fez++ H02' (2)
Fez++ H+ + O2
(3)
Ferric systems are attractive because degradation of organics can be catalytic in iron. It is now known that the oxidizing power of the Fenton-type systems can be greatly enhanced by irradiation with UV or UV/visible light (3-6). The herbicides 2,4dichlorophenoxyacetic acid (2,4-D) and 2,4,5-trichlorophenoxyacetic acid (2,4,5-T) are completely mineralized to COz and HC1 by Fe3+/Hz02accompanied by Pyrexfiltered visible fluorescent light (3). The stoichiometry of 2,4-D mineralization was found to be 7 mol of Hz02and 4 mol of O2 per mole of herbicide. The dark (thermal) Fe3+/H202mineralization is markedly slower and does not 304
Environ. Sci. Technol., Vol. 27, No. 2, 1993
go to completion, even at high [H,O,]. The generation of OH' in Fe3+/H,02/UVsystems was suggested by kinetic experiments (4). Although HzOzphotolyzes to give OH", it absorbs weakly at wavelengths above the Pyrex cutoff, and we confirmed the negligible role of direct peroxide photolysis. At shorter wavelengths (e.g., 254 nm), where peroxide photolyzes more efficiently, hydroxylation of 2-hydroxybenzoic acid by H20z/hvwas found also to be accelerated by Fe(II1) (5). This paper presents a detailed investigation of 2,4-D mineralization by Fe3+/H2OZ/hv,focusing on the photochemical reactions that contribute to enhancement. It seemed likely that at least one contributing reaction to photoenhancement is photolysis of aquated ferric ion [hereafter, Fe(III), when only HzO or OH- ligands are present]:
-
Fe(III)aq
hu
+
Fe2+ OH'
(4)
The photoreduced iron may be a precursor to a second OH' when peroxide is present via the Fenton reaction (eq 1). Several Fe(III), species are known to be photoactive, including Fe3+, $eOH2+, and Fe2(OH)z4+(water ligands omitted), and their importance depends on pH and wavelength of the emitted light (7). Photolysis occurs by dissociation of the ligand-to-metal charge-transfer (LMCT) exci$d-state complex. Luiitlk and co-workers (5) attributed Fe(II1)-enhanced HZO2/hv hydroxylation of 2hydroxybenzoic acid to reaction 4. Larson et al. (8) exploited reaction 4 for transformation of s-triazine herbicides in water by Fe3+/hv (no peroxide). Our original study indicated, however, that other reactions besides photolysis of Fe(III), are involved in Fe3+/HZOz/hvsystems. We show here that a number of reactions involving Fe(II1) species contribute. It should be noted that this chemistry is relevant to the environment, as well. Skurlatov et al. (9) suggested that iron with or without peroxide may contribute to sunlight photolysis of 2,4,5-T in natural waters. Also, iron ions and peroxide are believed to be involved directly or indirectly in photochemical oxidations in cloud droplets (7, 10, 11). Experimental Section Materials. The herbicide 2,4-D and Fe(C104)2.6Hz0 were from Aldrich. Anhydrous Fe(C104)3and Ce(HS04)4 were from GFS Chemicals (Columbus, OH). [ring-ULI4C]- and [c~rboxy-~~C]-2,4-D were from Sigma. All chemicals were used as purchased. General Procedures. Procedures are summarized except where they differ appreciably from earlier descriptions (3). Reaction solutions, typically 100 mL in a 250-mL Erlenmeyer flask, were stirred magnetically to keep the solution well-aerated. Iron solutions were made fresh. Reactions were initiated by addition of Hz02,and samples were quenched with methanol prior to analysis. A syringe pump (Model 355, Sage Instruments) was employed when continuous addition of HP02was required. The following conditions were used frequently and shall be referred to in tables, figures, and text as the "standard
0013-936X/93/0927-0304$04.00/0
0 1993 American Chemical Society
conditions": [Fe3+]= 1.0 mM, [2,4-D] = 0.1 mM, [HZ021 = 10 mM, pH 2.8, T = 25 OC, ionic strength 0.2 M (NaC104), UV light intensity 9 X lo1' quanta L-' s-l. Any deviations from these will be noted. For reactions under N2 or O2 atmosphere, the reaction flask was fitted with a serum stopper, which allowed entry of needles for purging and sampling. Solutions were bubbled with gas during the reaction, as well as for at least 0.5 h before initiation. Photolysis. The light source was four 15-W fluorescent black light blue tubes (F15T8/BLB, General Electric), which emit in the range 300-400 nm, centered around 360 nm (6). Light intensity was controlled by adjusting the distance from the reaction vessel and measured by ferri(12). oxalate actinometry [65 mL of 0.006 M K3[Fe(Cz04)311 Most experiments were carried out at 9.0 X 1017quanta L-l s-l (or -2.15 mW/cm2 at the solution surface, as measured with a quantum radiometer), with air in the headspace, unless specified otherwise. In cases where dioxygen was present, it was never limiting. Photoassisted H202decomposition by Fe3+ was conducted using a parabolic tungsten lamp filtered through CuS04 and NaNO, chemical filters, which allows transmitance between about 410 and 500 nM (12). Reaction solutions were presaturated with oxygen, and oxygen evolution was measured manometrically. Analyses. 2,4-D was determined by reverse-phase liquid chromatography. Mineralization was monitored by the loss of solution radioactivity from [ring-UL-14C]- or [~arboxy-~~C]-2,4-D. Loss of solution radioactivity was shown previously to correspond to appearance of 14C02in the headspace trap (3). Radioactivity was measured by liquid scintillation. UV/visible absorption spectra were recorded on a double-beam Perkin-Elmer spectrophotometer (Hitachi 200) at 60 nm/min. Oxalic acid was identified as its di-n-propyl ester derivative by gas chromatography/mass spectrometry (Hewlett-Packard 5890 GC/5970 MSD) using a DB-1 capillary column (J&W Associates). To obtain the derivative, the reaction mixture was neutralized, filtered to remove ferric oxyhydroxides, and freeze-dried. The residue was treated with 7.5% BF, in n-propanol at 100 "C for 2 min, followed by 5-fold dilution with water and extraction of the ester into chloroform (13). Authentic oxalic acid derivatized the same way gave an identical mass spectrum. Rate Simulations. The multistep kinetic program GEAR PC Version 1.11(Project Seraphim, Department of Chemistry, University of Wisconsin, Madison, WI 53706), based on the original Fortran HAVCHM (14),was used to simulate the kinetics of 2,4-D loss. The following rate constanta (M-l s-') were taken from the literature: kl = 53 (15);k2 = 0.02 extrapolated from Walling and Goosen (16) by assuming it to be proportional to [H+]-l; k-, = 1.2 X lo6 (17);the ratio k3lk-2 = 0.024 (16). Additional reactions which must be included are as follows:
--.
+ H202 HzO + HOP' k, = 2.7 X lo7 (18) OH' + Fe2+ OH- + Fe3+ k6 = 4.3 X lo8 (18) OH' + RH R' or [HORH]' (R = 2,4-D) OH'
-
k7 = 2.7 R' -t 0
2
RO2'
R02'
k, = 1.7 X
X
(5)
(6) (7)
lo9 (estimate) (3)
-
k8
+ Fez+
= -(5-50) X108 (19)
R02-
+ Fe3+
(8)
(9)
lo6 [estimate based on R = (CHJ&(OH)I (19)
r 1 0 1M MeOH
.- _- - - - - -
m d .e
u
2
40 -
20
------__-___
~
- - -1: 1 =I -7 --- - - - -. Peroxide added:
-
-.
lOmM continuoualy - l O m M all at once - - 50mM atepaise . - - 50mM all at once - - 10mMandFe2+suppl at arrow
0
Results and Discussion Two-Stage Nature of the Reaction and Involvement of Hydroxyl Radical. As reported before (3),Fe3+/HzO2 in the dark dechlorinates 2,4-D and oxidizes about 40-4590 of the ring and carboxy carbons to C02. In the full system (Fe3+/H202/hv),this point marks a convenient dividing line between two stages of the total oxidation of the compound. It will be shown here that (i) the first stage is dominated by OH' reactions, and light contributes to production of OH', and (ii) the second stage (i.e., last -60% of carbon mineralization) occurs without OH' and is entirely the result of photochemical reactions. Firsbstage mineralization in the dark is strongly inhibited by OH' scavengers such as methanol (Figure 1). We attribute oxidation during this stage to OH' generated by classical thermal Fenton-type chemistry, of which eqs 1-3 are the key steps. Figure 1 further shows that the remaining ring label is mineralized only with difficulty in the dark. Continuous or repeated injections of peroxide or supplementation with Fenton's reagent gave only slight additional decrease in solution radioactivity. Very high peroxide concentrations were moderately more effective; for example, 500 mM Hz02added all at once resulted in 69% of ring label evolved as 14C02(3). From these observations it is clear that OH' is ineffective in oxidizing the organic intermediates remaining after the first stage is complete, most likely due to lower reactivity of the presumed highly oxygenated intermediates toward OH' and competition for OH' by other solutes. The hydroperoxyl radical (HO,') produced by the well-known Ce(JY)/H202reaction (20) was incapable of releasing any CO, from the ring (data not shown), consistent with the generally poor reactivity of HOz' compared to OH' toward organic compounds (17, 18). In contrast to dark conditions, mineralization of the ring and carboxy carbons is rapid and complete under near-UV light irradiation (Figure 2). At the highest light intensity employed, all solution radioactivity is eliminated in -30 min. No attempt was made to maximize the rate. Both transformation and mineralization of 2,4-D are accelerated by light. The mineralization cuve in Figure 2 at the lowest intensity illustrates, again, the two-stage nature of the reaction. The results depicted in Figure 3 prove that OH' participates marginally, if at all, in second-stage mineralization. In this experiment, the reaction was carried out in the dark to completion of the first stage and then irradiated. Second-stage mineralization is seen to occur at the same rate in the presence or absence of high concentrations Environ. Scl. Technoi., Voi. 27, No. 2, 1993 305
Table I. Initial Rates of 2,4-D Transformation by Fe3+/Bv as a Function of OH' Savenger Concentration
[MeOH], mM
I
rate, lo4 M/min
[MeOH], mM
rate, lo4 M/min
0
6.9
0.1 0.5
5.3 3.1
3.0 10 100
0.70 0.10 0.08
1.0
2.1 in 0.t Y methanal
100
1
0
3
2
4
80
n
Time. h
I
Figure 2. Effect of light intensity on 2,4D mineralization. Inset Is for [carboxy-"C]-2,4-D at [H202] = 1.0 mM. Otherwise, standard conditions apply.
2
60
4
d
40 W
R
A 0.lM t-&OH,
Qht
I 2
0
6
4
B
10
2.0
2.6
Time, b
a,
carbory- 14C 2.4-D 1.2
%
b
*
a
. .
a
A
*
A
with methanol: A
0
16.0
0.5
16.6
17.0
17.6
D
Fe/W
18.0
Time, h Flgure 3. First stage dark followed by second-stage photolytic minerallzatlon of 2,4-D showing the absence of any effect by OH' scavengers during the second stage. [MeOH] = [BuOH] = 0.1 M. Otherwise, standard conditions apply.
0.0
0.6
1.o
1.5
2,4-D half-lives Figure 5. (A, upper) Dark or UV-assisted transformation of 2,443 with or without OH' scavenger methanol. (B, lower) Comparison of carboxy-I4C decarboxylatlon ratios in OW-scavenged vs unscavenged reactions. Otherwise, standard conditions apply.
0.01
0.1
1
10
100
IH2023, mM
Figure 4. Initial rates of 2,4D disappearance as a function of [H202] by the full system, the dark system, and the photolyzed system in the absence of peroxide (otherwise standard conditions). Dashed curve is the simulated dark reaction. Dotted curve is the simulated full system rate, assuming only photoreduction of ferric ion (eq 9) and photodecarboxylation of ferric-2,eD complex (eq 10) contribute. Observed rates for H2021hu(not shown) were 10.3 pmol L-' min-'.
(0.1M) of OH' radical scavengers methanol or tert-butyl alcohol (21). Photochemical Reactions Participating in the First Stage. To sort out the various photolytic reactions, we compared initial rates of 2,4-D transformation by Fe3+/ H202/huand the component systems, Fe3+/H202,Fe3+/hu, and H20z/hv,over a range of [H,02] (Figure 4). Transformation rather than mineralization was chosen because it should be uncomplicated by contributions from the 306
Environ. Sci. Technoi., Vol. 27, No. 2, 1993
second stage. Initial rates (Le., first 10-2070 reaction), as opposed to derived rate constants, eliminate product effects and circumvent misinterpretation that could arise from differences in kinetic order among the systems. Rates by H20z/huwere negligibly slow, as observed when visible light was used (3), and will not be further discussed. We believe three photolytic reactions contribute to UV assistance. Two of these are revealed in the Fe3+f hv (no peroxide) reaction. This reaction has two componentsone that is strongly inhibited by OH' scavengers (methanol) and one that is not inhibited (Figure 5A, Table I). The inhibited component is obviously due to LMCT photolysis of Fe(III)aqto OH', as in eq 4. It is well-known that such photolysis leads to oxidation of organic compounds (e.g., alcohols) (21). Under our conditions (pH 2.8, no complexing ligands present), the first ferric ion hydrolysis product FeOH2+ is the most photoactive (7). FeOH2+ comprises -50% of total iron (3) and has an absorbance tail that extends well above 300 nm (7).The OH' quantum yield (eq 10) is 0.14 at 313 nm and 0.017 at 360 nm (7). FeOH2+2Fez++ OH'
(10) Reported rates of reaction 10 are fully consistent with the supposition that it plays an important role in our reactions. Faust and Hoigne (7) reported a half-life of -32 min for
tert-butyl alcohol-scavenged photolysis of monomeric Fe(III), at pH 4.0 irradiated with 313-nm monochromatic light at 5.6 X 10l6quant L-' s-l, and -20 min for clear-sky solar noon sunlight photolysis. Under the latter condition, sunlight emits UV at about the same order of magnitude intensity as our light source in standard experiments (-6 vs 2.1 mW/cm2, respectively). 2,4-D has a half-life of -55 min in Fe3+/hvreactions (see, for example, Figure 8, vide infra). The noninhibited reaction in Fe3+/hv systems is attributable to LMCT photodecarboxylation of a 2,4-DFe(II1) complex: ArOCH,C02H
e
+ Fe3+
[Ar0CH2CO2-Fel2+
hu
[ArOCH,'] + CO, + Fe2+ (11) Photolysis of ferric carboxylate complexes (or their ion pairs) by a LMCT pathway is fairly common (22),resulting typically in reduction of iron and evolution of CO,. The organoradical is susceptible to further oxidation by 02. One good example is ferric trioxalate, which is widely used as a chemical actinometer (12). Zepp et al. (11) photolyzed ferric citrate and oxalate complexes to generate Fe2+in situ for Fenton reactions. We found no UV/visible spectral evidence for [ArOCH2CO2-FeI2+,but this does not preclude its formation in small amounts. Reaction 11 is supported by results in Figure 5 and Table I. Table I shows that the scavenging ability of methanol in Fe3+/hvsystems increases rapidly with concentration to 10 mM and then abruptly reaches a plateau, revealing the much slower, unscavenged reaction. Figure 5A depicts UV acceleration of OH'-scavenged reactions, with or without peroxide. Figure 5B shows that 14C02from carboxy-labeled 2,4-D is evolved in 1:l ratio with 2,4-D transformation in OH*-scavengedsystems (Fe3+/H202/hv and Fe3+/hv),but only in 0.2:l ratio in reactions without scavenger. Thus, in scavenged reactions a photoinduced decarboxylation predominates, presumably by eq 11,while in unscavenged systems O H reacts with parent compound at least -75% of the time in ways that do not lead to decarboxylation, such as attack on the aromatic ring. Whether photodecarboxylation proceeds in one step, as written in eq 11, or by a multistep pathway, as observed for certain polydentate ferricarboxylates (e.g., oxalate and tartarate) (22),remains undetermined. Regardless, the net result is decarboxylation simultaneous with 2,4-D loss. Also consistent with eq 11 as written is the nearly stoichiometric formation in methanol-inhibited Fe3+/hv reactions of 2,4-dichlorophenol (data not shown), which along with formate is the expected product from ArOCH; after further oxidation and hydrolysis. In unscavenged systems, considerably less (ca. 5-20%) of the phenol is detected. Inspection of Figure 4 reveals that photoreduction and photodecarboxylation (eqs 10 and 11)cannot completely account for UV acceleration of 2,4-D transformation. The effect of these reactions is to increase the steady-state concentration of OH'. Since Fez+ is a product, the rate in the presence of peroxide is potentially almost double the rate observed in its absence. A numerical multistep integration program GEAR was used to predict dark reaction rates (dotted curve) using literature or estimated rate constants for eqs 1-3 and 5-9. Next, the full system reaction rates were predicted (dashed curve) by including the measured rate constant of the Fe3+/hvreaction as the combined rate constant for eqs 10 and 11,and assuming that no other photoreaction occurs. (The rate of eq 11is small compared with eq 10.) Two things are apparent from the predictions. First, reaction 10 does not meet its potential doubling effect, most likely due to competitive
a
- mixture
0
6
10
15
20
25
Time, min Figure 6. Dark and photoassisted decomposition of H202with 410500-nm wavelength light, 5.6 X 10'' quanta L-' s-'; [H202] = 9 M, [Fe3+] = 1 mM, pH 2.8. Oxygen evolution was nonlinear during the first 3 mln (omitted) in the dark, possibly due to some siugglsh equilibrium. Inset: UV/vislble spectra of 4.5 M H,02, 0.5 mM Fe(III)aq, and a mixture of the two, all at pH 2 8. Employing equilibrium constants for hydrolysis of Fe3+(7) and formation of Fe-02H2+ (K,) (23) affords = 0.78. The calculated log t (extinction the ratio [Fe-02H2+]/[Fe],, coefficient) at 440 nm is 2.3, in agreement with the reported value of 2.6 (23).
reaction of Fe2+ with H02* (eq -2b) and RO,' (eq 9). Second, and most importantly, the observed full-system rates substantially exceed the rates predicted from just eqs 10 and 11 over most of the range in [H,O,]. These results suggest an additional Fe(II1)-sensitized photoreaction. One possibility is photolysis of the monoperoxy complex of Fe(III), which is formed in the preequilibrium step, eq 2a:
hv
Fe-02H2+
?
(12)
The formation constant of Fe-02H2+,K,, determined from We observed UV/visible spectral data (23)is 3.6 X similar spectral changes (Figure 6 inset). Note that the absorbance of Fe-02H2+ is greater, and extends further into the visible region, than Fe(III)aq,which is mainly that of FeOH2+above 300 nm. Under our standard conditions, when [H202]= 10 mM, Fe-02H2+ comprises initially 0.7% of total Fe(II1). Reaction 12, with Fez+and HO,' as products (presumably from LMCT excitation), was proposed by Behar and Stein (24),who observed temperature-independent acceleration of the Fe(III),-catdyzed decomposition of H20z to 0, at 365 nm. Their conclusions are suspect since both FeOH2+and H,02 absorb at 365 nm, albeit weakly, and they too are expected to give O2 when photolyzed, via eqs 5 and 3. Figure 6 shows, however, that irradiation above 410 nm, which is above the absorption edge of FeOH2+and H20z,increases the rate of 0,evolution -2-fold compared to the dark reaction. The rate attributable to photolysis was identical (within 20%) at 1and 25 'c, consistent with a photochemical reaction. These results support but do not prove reaction 12 (especially with regard to the photoproducts), and consequently the identity of the additional contributing reaction(s) remains speculative at this time. Unfortunately, the 0,evolution rates in Figure 6 cannot be meaningfully linked to 2,4-D degradation rates because (i) the stoichiometry between the oxidant species formed and O2 evolved is unknown, particularly since O2 may result from chain reactions (24),and (ii) the quantum yields in the visible region may be much smaller than in the UV. Photolytic Reactions Occurring in the Second Stage. To identify the required reactants in second-stage mineralization, the following experiment was conducted
-
Environ. Sci. Technoi., Vol. 27, No. 2, 1993 307
100
At
WOW:
EDTA added Appt. Fa. fine-filtered
!R5G-
Under N2
I
80
he
d
-
3 a m
.a
0
u
2
5
20 0
0
16
0.5
17
18
50
0
150
100
200
250
Time, h
Flgure 7. First stage dark followed by second-stage photolytic mineralization of 2,eD showing the Influence of Fe(II1) during the second stage. “EDTA added” means 10 mM ethylenediiminetetraacetlc acM was added prlor to photolysis. “Fine-flRered” means removal of colloidal ferric oxyhydroxlde precipitates by 0 . 4 5 - ~ mfilter, and “coarsaflltered”means removal of the same by coarse filter paper. Standard condltlons apply.
(Figure 7). The dark reaction was carried out for 16 h, giving the normal -40% loss of solution 14C. By this time, [H,O,] was completely decomposed, partly due to its consumption in first-stage reactions, and partly to Fe3+catalyzed decomposition, which has a reaction half-life of -1 h under those conditions (3). Two of the replicates were then made alkaline to pH 10 with NaOH to precipitate ferric oxyhydroxide, which was then removed by either fine or coarse filtration, and the filtrate was reacidified with HC104to the original pH of 2.8. Note that precipitation/filtration/reacidificationleft the solutionphase radioactivity virtually unchanged. A third replicate was untreated. UV irradiation of the untreated replicate resulted in rapid and quantitative volatization of the remaining label, whereas the fine-filtered replicate was almost unaffected. The coarse-filtered sample photolyzed slowly, apparently because a fraction of the colloidal iron oxyhydroxide passed through the filter paper and dissolved on reacidification. Finally, addition of an iron chelator, EDTA, to a fourth replicate completely eliminated photolytic volatization of label. These results demonstrate that second-stage mineralization (i) is due to ferric-sensitized reactions, (ii) does not require peroxide, and (iii) does not involve direct photolysis of the organics. Recall that OH’ is not involved in this stage. We attribute the ferric-sensitized reactions to LMCT photolysis of Fe(II1) complexes of degradation intermediates (eqs 13 and 14). These intermediates are most [RC0,-FeI2+
hv
-
[ROH, FeI3+
hv
[R’] + C 0 2 + Fez+
(13)
[RO]’ + Fez+ + H+
(14)
likely polydentate ligands containing carboxyl and hydroxyl groups. The absorption spectrum of reaction solutions taken at the end of first stage showed tailing into the visible region beyond that of Fe(III), (3),consistent with complexation. Oxalic acid was identiaed as a product in reaction mixtures after the dark reaction was complete, and it, of course, is photolabile well into the visible region. Photodissocation of the Fe(II1)-carboxylate bond (eq 13) is well-established, as discussed above. Photodissociation of RO-Fe(II1) complexes (eq 13) was shown to take place in neat alcohol solutions (22,25). Phenol reportedly photolyzes at wavelengths of >400 nm in dilute solution containing ferric ion (26). Ethylene glycol (1 M) photolyzes 308
Environ. Sci. Technol., Vol. 27, No. 2, 1993
0
OmMH202
0
0.1
* 1 A
100
60
0
10
160
200
250
Time, min Flgure 8. 2,4-D mineralization in N, (A, upper) or 0,(B, lower) at various [H20,]. Otherwise, standard condltlons apply.
to formaldehyde and Fez+ with a quantum yield of 0.05 at 254 nm (27). It is clear from Figure 7 together with data showing no OH’ contribution (Figures 1 and 3) that reactions like eqs 13 and 14 are solely responsible for second-stage mineralization. Role of Oxygen in the Reaction. As mentioned previously (3), dioxygen plays an important role in these reactions. Dioxygen is consumed, at least when [H20,] is relatively low. Oxygen incorporation is known to occur in the oxidation of benzene by OH’/*802(28). Figure 8 shows that reactions under N2 are slower and require much more peroxide to go to completion than reactions under an 0, headspace. (Reactions in air or 0, headspace gave identical results.) This is true even for initial rates. Notably, initial rates in the absence of H202(topmost curves) are enhanced by roughly a factor of 4. Dioxygen enhancement can be attributed to the following mechanisms, which may be operative under different peroxide regimes or at different times during the reaction. 1. Dioxygen reacts with intermediate organoradicals to give products containing hydroxyl and carboxyl functional groups, which then form photolabile Fe(II1) complexes, thereby promoting overall photomineralization. 2. Dioxygen reacts with intermediate organoradicals to regenerate H202by the “Dorfman” mechanism (eqs 15 and 16) (29). This is especially important at low added [H20,1, R’
+ 02
-+
ROz’
HZO
H+ + H02’ + [e-]
-
ROH
+ HOz’
HzOz
(15) (16)
where the reaction may be quite sensitive to peroxide concentration. Electrons in eq 16 may come from organoradicals or Fez+. 3. Intermediate hydroperoxy, organoperoxy, and other organoradicals formed by reaction with O2 serve to reoxidize the Fez+ photoproduct to Fe(II1) (eq 9), which is
Scheme I. Pathway of OH' Generation (A) and Pathways of Chlorophenoxy Herbicide Mineralization (B)"
Table 11. Mineralization of 2,4-D by Fe2+/hv" % initial
time, min 0
10 20
30 60 90 120
I4C remaining
ring-labeled
carboxy-labeled
100 99.6 98.6 96.4 90.9 83.8 67.7
100 96.0 95.6 94.6 81.0 56.2 40.3
ring-labeled plus 10 mM HzOz 100 25.5 7.7 3.1 0.1
ll Fe3
+
Standard conditions apply.
photoactive. This process is vital only during the second stage. [Direct oxidation of Fe2+by O2 (autoxidation) is too slow in acidic solution to be important (3, 30).] Finally, Table I1 shows that irradiated, oxygenated ferrous systems (Fe2+/hv) are capable of mineralizing 2,4-D, but only after a fairly long lag. When peroxide is added the lag disappears because the Fe2+is immediately oxidized to Fe3+. The lag could be due to slow autoxidation of Fez+to generate small amounts of Fe3+,which may then photocatalyze degradation via eq 10, etc. Each organoradical produced can oxidize potentially two Fe2+via eq 9, eq -2b, and/or the reaction analogous to eq 1involving organoperoxide instead of hydrogen peroxide; Le., a net chain-carrying process as 2Fe2++ O2 R' 2Fe3+ + RO'(0H') HO-(RO-) (17) Thus Fe3+generation can be autocatalytic. In any case, the long lag clearly indicates that Fe2+photochemical reactions (22) are unimportant to the rest of the chemistry described in this paper regarding 2,4-D mineralization.
+
-
6
co2 hV
[ArOCH2C02Fe12
Fe3
(side-chain attack
HCI t C 0 2
OH(ring
attack)
HCI + C 0 2
+
Summary and Mechanistic Conclusions The pathway for total mineralization of 2,4-D by Fe3+/H202/hvis summarized in Scheme I. Part A shows the steps leading to OH'. The thermal reaction and photolysis of FeOH2+generate O H directly. All photolytic reactions may contribute indirectly to generation of OH' via production of Fez+by the Fenton reaction. The importance of individual reactions depends on [H202]and, obviously, on light intensity. It can be seen from Figure 4 that photoreduction of ferric ion is prominant at low peroxide, the dark reaction at high peroxide, and the unassigned photoreaction(s) at intermediate peroxide levels. In part B, the relative importance of the pathways is indicated by the thickness of the arrow. A small fraction, at most 1%,of 2,4-D is lost by charge-transfer photodecarboxylationof a ferric-2,4-D complex. About 2@-25% of net OH' reactions with 2,4-D are attacks on the side chain leading to simultaneous decarboxylation. In agreement with this is the consistent production of about 8-16 9% transient yield of 2,4-dichlorophenol, the expected product from attack on either the carboxyl or CH2groups. The other -75% of OH' attacks are probably on the aromatic ring. Oxidation of the ring by OH'/02 causes dechlorination and partial mineralization of ring carbon. The eventual ring-opened products, one of which was identified aa oxalate, form complexes with Fe(II1) that are photolyzed, leading to further evolution of C02 and probably consumption of OD Hydroxyl plays little role in mineralization of -60% of the ring carbon. Oxygen is consumed in the reaction and accelerates degradation by producing H202,oxygenated intermediates suitable for charge-transfer photolysis as Fe(II1) complexes, and oxy-
'
'
coi [ ROFe]
hv
R,.RO'
4
co2
Width of arrows denotes relative importance.
genated organoradicals that maintain Fe(II1) concentrations during later stages of degradation. We are presently looking at the decomposition of other organic compounds by Fe3+/H202/UVto evaluate this system for treatment of hazardous chemicals. Literature Cited (1) Peyton, G. R. In Emerging Technologies in Hazardous
(2) (3) (4)
(5) (6)
(7) (8) (9)
Waste Management;Tedder, D. W., Pohland, F. G., Eds.; ACS Symposium Series 422; American Chemical Society: Washington, DC, 1990; pp 100-118. O h ,D. F.; Pelizzetti, E.; Serpone, N. Enuiron. Sci. Technol. 1991,25, 1523-1529. Pignatello, J. J. Enuiron. Sci. Technol. 1992,26,944-951. Haag,W. R.; Yao, C. C. D. Enuiron Sci. Technol. 1992,26, 1005-1013. SedlBk, P.; LufiBk, S.; Lederer, P. Collect. Czech. Chem. Commun. 1987,52,2451-2456. Pignatello, J. J.; Sun, Y. In Emerging Technologies in Hazardous Waste Management;Tedder, D. W., Pohland, F. G., Eds.; American Chemical Society: Washington, DC, Vol. 111, in press. Faust, B. C.; HoignB, J. Atmos. Enuiron. 1990,24A, 79-89. Larson, R. A.; Schlauch, M. B.; Marley, K. A. J. Agric. Food Chem. 1991,39, 2057-2062. Skurhtov, Y. I.; Zepp, R. G.;Baughman, G. L. J. Agric. Food Chem. 1983,31, 1065-1071. Environ. Sci. Technol., Voi. 27, No. 2, 1993 909
Environ. Sci. Technol. 1993,27,310-315 Hoign6, J.; Zuo, Y.; Nowell, L. Preprints of Extended Abstracts, 203rd National Meeting of the American Chemical Society, Division of Environmental Chemistry, San Francisco, CA; American Chemical Society: Washington, DC, 1992; Paper 92, pp 147-149. Zepp, R. G.; Faust, B. C.; Hoign6, J. Environ. Sei. Technol. 1992, 26, 313-319. Calvert, J. G.; Pitts, J. N., Jr. Photochemistry; John Wiley and Sons: New York, 1966; pp 737-786. Appleby, A. J.; Mayne, J. E. 0. J. Gas Chromotogr. 1967, 5, 266-268. Stabler, R. N.; Chesnick, J. Int. J . Chem. Kinet. 1978, I O , 461-469. Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R. Trans. Faraday SOC. 1951, 47, 462-500. Walling, C.; Goosen, A. J. Am. Chem. Soc. 1973, 95, 2987-2991. Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L. J. Phys. Chem. Ref. Data 1985, 14, 1041-1100. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988,17, 513-886. Neta, P.; Hule, R. E.; Ross, A. B. J. Phys. Chem. Ref. Data 1990, 19, 413-513. Metelitsa, D. I. Russ. Chem. Rev. (Engl. Transl.) 1971,40, 563-580.
Langford, C. H.; Carey, J. H. Can. J. Chem. 1975, 53, 2430-2435. Balzani, V.; Carassiti, V. Photochemistry of Coordination Compounds; Academic Press: London, 1970; Chapter 10, pp 145-192. Evans, M. G.; George, P.; Uri, N. Trans. Faraday SOC.1949, 44, 230-236. Behar, B.; Stein, G. Science 1966, 154, 1012. Walling, C. H.; Humphreys, W. R. J. Org. Chem. 1981,46, 1260-1263. Dufek, P.; Cernohorsky, I.; PacBkovB, V. J. Chromatogr. 1982,241, 19-28. Carey, J. H.; Cosgrove, E. G.; Oliver, B. G. Can. J. Chem. 1977, 55, 625-629. Kunai, A.; Hata, S.; Ito, S.; Sasaki, K. J. Am. Chem. Soc. 1986, 108, 6012-6016. Dorfman, L. M.; Taub, I. A.; Buhler, R. E. J. Chem. Phys. 1962,36, 3051. George, P. Trans. Faraday Soc. 1954,50,4349-4359.
Received for review M a y 28, 1992. Revised manuscript received October 5, 1992. Accepted October 19, 1992. Support was received f r o m US.Department of Agriculture Pesticide Impact and Water Quality Special Grants Programs.
Reduction of Trace Element Concentrations in Alkaline Waste Porewaters by Dedolomitkation Eric J. Reardon,” C. James Warren, and Monlque Y. Hobbs
Department of Earth Sciences, University of Waterloo, Waterloo, Ontario, Canada, N2L 3G1 Dolomite [CaMg(C03)2]addition to alkaline waste materials, such as high lime content fly ash, is proposed as a method to reduce the concentrations of undesirable elements in leachate waters. The results of this study indicate that dolomite, in the presence of portlandite [Ca(OH)2],undergoes conversion to an assemblage of brucite [Mg(OH),] and calcite (CaCOJ in unstirred pastes reacted for 8 months. Calcite coated the surfaces of reacted dolomite fragments while brucite occurred as individual acicular crystals. The generation of calcite can reduce porewater concentrations of certain trace elements through coprecipitation mechanisms. The results of experiments conducted in this study showed moderate uptake of B and Se by precipitating calcite. Arsenic was included in this evaluation, but the high pH of the reaction solution resulted in its immediate precipitation as basic calcium arsenate [Ca4(As04)2(OH)z.4Hz0]. The potential of dedolomitization to reduce the concentrations of other undesirable elements in alkaline waste porewaters needs to be investigated. Dolomite is readily available in many areas of North America and can be easily processed to yield a fineness that is reactive when mixed with alkaline waste material. 1. Introduction
Fly ash, a byproduct of coal-fired electrical generating stations, poses a major disposal problem. With the placement of more stringent limits by many countries on the trace element content and leachability of materials used as lakefill, landfill, and construction material, it appears that ever-increasing volumes of fly ash will be relegated to monitored disposal sites. Contamination of local groundwaters by fly ash leachates at these sites is of environmental concern. Although ash leaching characteristics 310
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vary widely, the elements most often observed to be above drinking water standards include B, As, Se, Mo, Cd, Cr, Zn, and S as SO4 (1). The precipitation of calcite (CaC03)may be a means to reduce trace element concentrations in fly ash leachates. Most natural groundwaters are at, or near, calcite saturation. If calcite could be induced to precipitate in fly ash waste and entrain undesirable trace elements through the process of coprecipitation, then these elements would be relatively secure from rerelease because of the unaggressiveness of background groundwaters to dissolve calcite. Calcite precipitation in fly ash can occur, or be induced to occur, by various mechanisms. Calcite may precipitate through reaction with natural sources of carbon dioxide due to the intrinsic carbonation capacity of fly ash. If the carbonation capacity of the ash is very low, it could be artificially increased through the addition of lime or cement. Some flue gas desulfurization technology involves the addition of slurried limestone into the stack, which would generate high lime contents in the ash, and thus high carbonation capacities (2). If trace element concentrations in ash leachates can be reduced by coprecipitation with calcite, a practical method should be developed to induce carbonate precipitation at ash disposal sites. Below the water table, influx of background groundwater would contribute carbonate. However, influx of groundwater into fly ash would simply displace the existing trace element-enriched porewater into the surrounding aquifer. Carbonation by background groundwater could encapsulate elements, such as Se, which are not heavily partitioned into the aqueous phase to begin with, but readily soluble elements such as B would escape encapsulation. Above the water table, fly ash carbonation would occur as a result of diffusion of atmospheric or soil
0013-936X/93/0927-0310$04.00/0
0 1993 American Chemical Society
